Electrolytic Cells, Corrosion and Batteries

Some things never change: oxidation is still at the anode, reduction is still at the cathode, and electrons still travel through the wire from anode to cathode. What changes is who does the pushing. The power supply pulls electrons out of the anode (making it positive) and pushes them onto the cathode (making it negative). The supply must provide at least as many volts as the size of the negative cell potential, and in practice somewhat more.

The salt must be molten (melted), not dissolved. In solid NaCl the ions cannot move, so no current flows. In aqueous NaCl, water is easier to reduce than ${\text{Na}}^{+}$ ($-0.83\text{ V}$ versus $-2.71\text{ V}$), so hydrogen gas forms at the cathode instead of sodium metal.

• Water. Pure water conducts poorly, so a small amount of an electrolyte such as sodium sulfate or sulfuric acid is added. Written for acidic solution, the half-reactions are $2{\text{H}}^{+}+2{\text{e}}^{-}\to {\text{H}}_2$ (0.00 V) at the cathode and $2{\text{H}}_2\text{O}\to {\text{O}}_2+4{\text{H}}^{+}+4{\text{e}}^{-}$ (reduction potential +1.23 V) at the anode. Overall $2{\text{H}}_2\text{O}\to 2{\text{H}}_2+{\text{O}}_2$, needing at least 1.23 V in theory and more in practice. Hydrogen and oxygen form in a 2 : 1 volume ratio.
• Aluminum. Aluminum oxide is dissolved in molten cryolite and electrolyzed; aluminum metal forms at the cathode, and the carbon anodes react with oxygen and are gradually used up. Aluminum is too reactive to be produced by reducing its ore with carbon, so electrolysis is essential.
• Zinc. Zinc is commonly produced by electrolysis of a purified zinc sulfate solution, with zinc depositing on the cathode.
• Copper refining. An impure copper anode and a pure copper cathode sit in copper(II) sulfate solution. Copper dissolves from the anode and plates onto the cathode. More active impurities (such as iron and zinc) dissolve and stay in solution; less active metals (such as silver and gold) do not oxidize and fall to the bottom as anode sludge.
• Electroplating. The object to be coated is made the cathode, the plating metal is the anode, and the electrolyte contains ions of the plating metal. Silver, copper and nickel plating commonly work this way. (Chromium plating is an exception: it uses an unreactive anode, and the chromium comes from compounds dissolved in the electrolyte.)
• Calculating amounts. Current $I$ (in amperes) flowing for time $t$ (in seconds) delivers charge $Q=It$ in coulombs. One mole of electrons carries about 96 485 C, the Faraday constant $F$. So $n({\text{e}}^{-})=\frac{It}{F}$, and the half-reaction's mole ratio converts moles of electrons into moles of metal deposited or gas produced.

Rusting iron is a tiny galvanic cell. Where a water droplet sits on steel, one region of the surface acts as an anode and iron is oxidized. Electrons travel through the metal itself to another region, usually near the edge of the droplet where oxygen is plentiful, which acts as the cathode. The ${\text{Fe}}^{2+}$ and ${\text{OH}}^{-}$ ions meet and form iron(II) hydroxide, which is further oxidized by air to rust, hydrated iron(III) oxide, ${\text{Fe}}_2{\text{O}}_3\cdot x{\text{H}}_2\text{O}$. Rust is flaky and porous, so it does not protect the metal underneath.

• Primary cells cannot be practically recharged, e.g. the alkaline cell (zinc anode, manganese dioxide cathode, potassium hydroxide electrolyte, about 1.5 V).
• Secondary (rechargeable) cells can be recharged by running the reaction backward as an electrolytic cell. A lead-acid car battery has six cells of about 2 V each for about 12 V: $\text{Pb}+{\text{PbO}}_2+2{\text{H}}_2{\text{SO}}_4\to 2{\text{PbSO}}_4+2{\text{H}}_2\text{O}$ during discharge. In a lithium-ion battery, lithium ions move between a graphite electrode and a metal oxide electrode as it charges and discharges.
• Fuel cells are galvanic cells fed continuously with fuel. A hydrogen fuel cell oxidizes ${\text{H}}_2$ at the anode and reduces ${\text{O}}_2$ at the cathode; the overall reaction is $2{\text{H}}_2+{\text{O}}_2\to 2{\text{H}}_2\text{O}$, with a standard cell potential of 1.23 V. The only product at the cell is water, but how clean the whole system is depends on how the hydrogen is made, and hydrogen is difficult to store and transport.
• Safe disposal. Batteries contain metals and electrolytes that harm ecosystems, such as lead, cadmium, nickel, lithium compounds and acids. Take used batteries to a battery recycling drop-off or household hazardous waste depot instead of the garbage. Damaged lithium-ion batteries can catch fire, so tape their terminals and never crush or puncture them.